Chapter 4 – Reactions in Aqueous Solution

4.1 – General Properties of Aqueous Solutions

An aqueous solution is a solution in which water is the solvent. Understanding the behavior of substances in water is fundamental to chemistry.

• Electrolyte: A substance that dissolves in water to produce ions, thus conducting electricity. Strong electrolytes dissociate completely (e.g., NaCl), while weak electrolytes dissociate partially (e.g., acetic acid).

• Nonelectrolyte: A substance that dissolves in water but does not produce ions (e.g., sugar).

• Examples: NaCl (strong electrolyte), CH3COOH (weak electrolyte), glucose (nonelectrolyte).

4.2 – Precipitation Reactions

Precipitation reactions occur when two solutions of ionic compounds are mixed and an insoluble product (precipitate) forms.

• Solubility chart: Used to determine if a compound is soluble (aq) or insoluble (s) in water.

• Precipitate: An insoluble solid formed in a reaction, usually denoted as (s).

• Ionic equation: Shows all ions present in a reaction. Net ionic equation: Shows only the ions and molecules directly involved in the reaction.

• Spectator ions: Ions that do not participate in the reaction and remain unchanged.

Example: Mixing AgNO3 and NaCl forms AgCl (s) as a precipitate.

4.3 – Acids, Bases, and Neutralization Reactions

Acids and bases are fundamental chemical species in aqueous reactions. Their strength and behavior are important for predicting reaction outcomes.

• Acid: Produces H+ ions in solution. Base: Produces OH- ions.

• Strong acids/bases: Completely dissociate in water. Weak acids/bases: Partially dissociate.

• Neutralization reaction: Acid reacts with base to form water and a salt.

• Common products: Water (H2O) and an ionic compound (salt).

• Net ionic equation for strong acid/base:

4.4 – Oxidation-Reduction Reactions

Redox reactions involve the transfer of electrons between species. Understanding oxidation numbers is key to identifying these reactions.

• Oxidation: Loss of electrons. Reduction: Gain of electrons. (OIL RIG: Oxidation Is Loss, Reduction Is Gain)

• Oxidation number: Assigned to elements to track electron transfer. Rules are based on element type and compound structure.

• Example: In NaCl, Na has +1, Cl has -1 oxidation number.

4.5 – Concentrations of Solutions

Concentration describes the amount of solute in a given volume of solution. Molarity is the most common unit.

• Molarity (M):

• Conversion: 1 L = 1000 mL

• Dilution equation:

• Ion concentration: For Na2SO4, [Na+] = 2 × [Na2SO4], [SO42-] = [Na2SO4]

4.6 – Solution Stoichiometry and Chemical Analysis

Stoichiometry in solutions involves using concentrations and volumes to calculate reactant and product quantities.

• Titration: A method to determine concentration by reacting a known volume with a solution of known concentration.

• Dimensional analysis: Used to convert units and solve stoichiometry problems.

Chapter 6 – Electronic Structure of Atoms

6.1-6.4 – Nature/Behavior of Light, Quantized Energy, and the Bohr Model

Light and energy are quantized in atomic systems. The Bohr model explains electron transitions and energy levels.

• Frequency (ν): Number of cycles per second (Hz).

• Wavelength (λ): Distance between peaks (meters).

• Relationship: (c = speed of light, m/s)

• Photon energy: (h = Planck's constant, J·s)

• Bohr model: Electrons occupy quantized energy levels; transitions emit/absorb photons.

• Energy change:

• Emission: Electron drops to lower energy level, emitting photon. Excitation: Electron absorbs energy and moves to higher level.

6.5/6.6 – Quantum Mechanics, Atomic Orbitals/Representations of Orbitals

Quantum numbers describe the properties and locations of electrons in atoms.

• Principal quantum number (n): Indicates energy level; n = 1, 2, 3, ...

• Angular momentum quantum number (l): Defines orbital shape; l = 0 (s), 1 (p), 2 (d), 3 (f)

• Magnetic quantum number (ml): Specifies orientation; ml = -l to +l

• Shapes: s (sphere), p (dumbbell), d (cloverleaf)

6.7-6.9 – Many-Electron Atoms, Orbital Diagrams, and Electron Configurations

Electron configurations and orbital diagrams show how electrons fill atomic orbitals according to specific rules.

• Spin quantum number (ms): +1/2 or -1/2; describes electron spin.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Electron configuration: Order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, ...

• Condensed configuration: Uses noble gas core to simplify notation (e.g., [Ne]3s23p1).

• Orbital capacity: s (2 electrons), p (6), d (10)

Example: Electron configuration for sodium (Na): 1s22s22p63s1; condensed: [Ne]3s1.